This will help with the basics of electrochemistry

1. Electrochemistry
Mr Aphane
Grade 12
Balancing Redox Reactions
Galvanic Cells
Standard Reduction Potentials
Electrolytic cell and its applications

2. Oxidation-Reduction (Redox) Reactions
• Review of Terms:
Oxidation–reduction (redox) reactions involves
transfer of electrons from one reactant (the
reducing agent) to another (the oxidizing agent)
Oxidation – the loss of electrons
Reduction – the gain of electrons
Reducing agent – electron donor
Oxidizing agent – electron acceptor

3. Redox Reactions
• Examples of Redox Reactions:
Zn(s) + CuSO4(aq)  ZnSO4(aq) + Cu(s);
Zn(s) + Cu2+
(aq)  Zn2+
(aq) + Cu(s)
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s);
Cu(s) + 2Ag+
(aq)  Cu2+
(aq) + 2Ag(s)
MnO4
‫־‬
(aq) + 5Fe2+
(aq) + 8H+
(aq) 
Mn2+
(aq) + 5Fe3+
(aq) + 4H2O

4. Balancing Redox Equations:
The Half-Reaction Method
1. Write separate equations for oxidation and
reduction half–reactions.
2. For each half–reaction:
Balance all the elements except H and O.
Balance O using H2O.
Balance H using H+.
Balance the charge using electrons.
3. If necessary, multiply one or both balanced half–
reactions by an integer to make the number of
electrons in both half–reactions equal.
4. Add half–reactions and cancel identical species.

5. Balancing Redox Equations
Example: balancing a redox reaction under acidic
condition
Cr2O7
2-(aq) + HSO3
-(aq)  Cr3+(aq) + HSO4
-(aq)
• How can we balance this equation?
• First Steps:
 Separate into half-reactions.
 Balance elements except H and O.

6. Balancing Redox Equations in Basic Solution
1. Use the half–reaction method as specified for acidic
solutions to obtain the final balanced equation as if
H+ ions were present.
2. To both sides of the equation, add a number of OH–
ions that is equal to the number of H+ ions present.
(You want to eliminate H+ by turning is into H2O)
3. Form H2O on the side containing both H+ and OH–
ions, and eliminate the number of H2O molecules
that appear on both sides of the equation.
4. Check that elements and charges are balanced.

7. Sample Exercises
• Balance the following redox reactions in basic
solution:
• Br2(aq) + OH-(aq)  BrO3
-(aq) + Br-(aq) + H2O;
• Cr(OH)4
-(aq) + OH-(aq)  CrO4
2-(aq) + H2O;

8. Applications of Redox Reactions
• Redox reactions such as combustion reactions are
very exothermic – they have very large negative DH;
• Redox reactions in aqueous solution also have
negative DH
• Available free energy from spontaneous reactions can
be trapped to produce electricity;
• Devices that utilize redox reactions to produce
electricity are called Galvanic cells or batteries.

9. Electrode Potentials and
Their Measurement
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
Cu(s) + Zn2+(aq)
No reaction

10. Terminology
• Galvanic cell:
A device that produces electricity from
spontaneous redox reactions.
• Electrolytic cell:
A device the uses electrical energy to make a
nonspontaneous chemical reaction to occur.

11. Galvanic Cell
• A device in which chemical energy is
converted to electrical energy.
• It uses a spontaneous redox reaction to
produce a current that can be used to generate
energy or to do work.

12. A Galvanic Cell

13. In Galvanic Cell:
• Oxidation occurs at the anode.
• Reduction occurs at the cathode.
• Salt bridge or porous disk allows ions to flow
without extensive mixing of the solutions.

14. Electrochemical Terminologies
• Anode half-cell - where oxidation process occurs;
• Cathode half-cell - where reduction process occurs;
• Electricity – electrons flow in the wire from the anode to
the cathode half-cells; in solution, cations and anions
flow in opposite directions across the salt bridge.
• Cell potential (Ecell) - electromotive force (emf) that
drives electrons and ions to flow; aka electrical potential.
 The unit of electrical potential is volt (V).
1 V = 1 J/C (Joule/Coulomb of charge transferred)

15. Standard Electrode Potentials
• Cell voltage: the electrical potential difference of
an electrode-pair.
• The cell potential of individual electrodes are
measured against the Standard Hydrogen
Electrode (SHE), which is reference electrode
assigned an electrical potential value of 0.00 V.

16. Standard Hydrogen Electrode
2 H+(a = 1) + 2 e-  H2(g, 1 bar) E° = 0 V
Pt|H2(g, 1 bar)|H+(a = 1)

17. Measuring Standard Reduction Potential
cathode cathode anode
anode

18. Reduction Couples
Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ?
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s); E°cell = 0.340 V
Standard cell potential: the potential difference of a
cell formed from two standard electrodes.
E°cell = E°cathode - E°anode
cathode
anode

19. Standard Cell Potential
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s); E°cell = 0.340 V
E°cell = E°cathode - E°anode
E°cell = E°Cu2+/Cu - E°H+/H2
0.340 V = E°Cu2+/Cu - 0 V
E°Cu2+/Cu = +0.340 V
H2(g, 1 atm) + Cu2+(1 M) → H+(1 M) + Cu(s); E°cell = 0.340 V

20. Standard Reduction Potentials
• Reduction potential, E°, for other electrons are
assigned positive (+) or negative (-) values,
depending on whether their reduction potential is
greater or smaller than the reduction potential of
Hydrogen electrode under standard condition.
 Standard condition implies an electrolyte
concentration of 1 M or gas pressure of 1 atm, and
the temperature is 25°C (or 298 K)

22. A Cu-Zn Galvanic Cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s); E°cell = 1.103 V

23. An Ag-Cu Electrochemical Cell

24. Cell Notations for Galvanic Cells
• A short-hand to describe electrochemical cells.
Anode half-cell on the left.
Cathode half-cell on the right.
Half-cells are separated by double vertical lines (||).
The concentration of each solutions is indicated in the
notation if known.
• Example: Mg(s)|Mg2+(aq)||Al3+(aq)|Al(s)
• Half-cell reactions:
 Mg  Mg2+ + 2e– (at anode)
 Al3+ + 3e–  Al (at cathode)

25. Galvanic Cell Notation
• Electron flows from the anode to cathode;
• Conventional current flows from cathode to anode;
• Positive ions flows into cathode half-cell, and
negative ions flows into anode half-cell via the “salt
bridge”.

26. Corrosion
• Corrosion is an electrochemical process in
which the metal is oxidized.
• To prevent corrosion, the metal must be
protected from being oxidized.

27. Corrosion of Iron

28. Corrosion Prevention
• Apply coating (such as paint or metal plating)
 Galvanizing (covering with zinc)
• Alloying that prevent the metal of interest
from being oxidized;
• Anodic protection – corrosion protection for
some metals by their oxide coating;
• Cathodic protection;
used to protects underground steel pipes from
corrosion.

29. Electrolysis
• A process that forces a current through a cell
to produce a chemical change for which the
cell potential is negative.

30. Commercial Electrolytic Processes
• Production of aluminum
• Purification of metals
• Metal plating
• Chloro-alkali industry

31. The Hall-Heroult Process for Al Production

32. Electroplating/Silver Plating a Spoon

33. Chloro-alkali